GCSE Chemistry specification checklist

FoundationEvery student

• 1.1Describe the atom as a positively charged nucleus surrounded by negatively charged electrons, with the nuclear radius far smaller than the atomic radius and almost all the mass in the nucleus.
• 1.2Recall the relative electrical charges and approximate relative masses of protons, neutrons and electrons, and that atoms have a radius of about 1 × 10⁻¹⁰ m.
• 1.3Use the atomic number (protons) and mass number (protons + neutrons) to work out the numbers of protons, neutrons and electrons in atoms and simple ions.
• 1.4Describe isotopes as atoms of the same element with the same number of protons but different numbers of neutrons, and explain why they have identical chemical properties.
• 1.5Describe how the model of the atom developed from indivisible spheres, through the discovery of the electron and the nuclear model from alpha-particle scattering, to the shell model and the discovery of the neutron.
• 1.6Write electronic structures in the 2,8,8 shell-filling order, and relate the number of outer-shell electrons to group number and the number of shells to period number.
• 1.7Describe how the periodic table developed, including Mendeleev leaving gaps for undiscovered elements, and explain why ordering by atomic number (once protons were known) fixed the earlier anomalies.
• 1.8Compare the characteristic physical and chemical properties of metals with non-metals and relate the difference to atomic structure and bonding.
• 1.9Describe the reactions of Group 1 metals with water, oxygen and chlorine, and explain why their reactivity increases down the group.
• 1.10Describe the reactions of Group 7 elements with metals and with other halide solutions, and explain why their reactivity decreases down the group.
• 1.11Describe the properties of Group 0 noble gases (monatomic, colourless, unreactive) and explain the trend in their boiling points down the group.
• 1.12Predict reactions and reactivity of elements from their positions in the periodic table, including displacement reactions between halogens.
• 1.13Recall and explain the general properties of the transition metals: hard and strong, high melting points, high densities, low reactivity compared with Group 1, ions with different charges, coloured compounds and uses as catalysts.Chemistry only
• 1.14Calculate the relative atomic mass of an element from the relative abundances of its isotopes.

FoundationEvery student

• 2.1Use the particle model to describe the arrangement, separation and movement of particles in solids, liquids and gases, and recall that the state at room temperature depends on the strength of the forces between particles.
• 2.2Describe melting, boiling, condensing and freezing as physical changes in which no new substance forms, and identify the energy transfers involved.
• 2.3Predict the physical state of a substance at a given temperature from its melting and boiling points.
• 2.4Explain, in terms of particles and the forces between them, why the temperature stays constant while a substance melts or boils even though it is still being heated.
• 2.6Distinguish a pure substance (a single element or compound) from a mixture, and use sharp, fixed melting and boiling points as evidence of purity.
• 2.7Describe a formulation as a mixture with a precisely designed composition and purpose, giving examples such as fuels, alloys, fertilisers, medicines and cleaning products.
• 2.8Choose and describe the separation technique suited to a given mixture: filtration for an insoluble solid, crystallisation for a soluble solid from solution, simple distillation for a solvent from a solution, and fractional distillation for miscible liquids with different boiling points.
• 2.9Describe paper chromatography as a separation of dissolved substances between a stationary phase and a mobile phase, and interpret simple chromatograms.
• 2.10Explain how fractional distillation separates crude oil and liquid air, relating each fraction to its boiling point range.
• 2.11Explain each separation choice in terms of the physical properties of the substances involved (solubility, boiling point, attraction to the stationary phase).

HigherHigher tier only

• 2.5Recognise the limitations of the particle model: it does not show the forces between particles, the sizes of particles relative to their separation, or the space between particles in solids and liquids accurately.

FoundationEvery student

• 3.1Identify when each strong bond forms: ionic bonding between a metal and a non-metal, covalent bonding between non-metals, and metallic bonding between metal atoms.
• 3.2Describe ionic bonding as electron transfer: metal atoms lose outer electrons to form positive ions, non-metal atoms gain electrons to form negative ions, and the oppositely charged ions are held by strong electrostatic forces of attraction.
• 3.3Draw dot-and-cross diagrams for ionic compounds formed from the first twenty elements, showing the charges on the ions and the electron arrangement of each ion, using square brackets.
• 3.4Deduce the formulae of ionic compounds from the charges on their ions, recalling common ion charges and compound ions such as OH⁻, NO₃⁻, SO₄²⁻ and CO₃²⁻.
• 3.5Describe the structure of an ionic compound as a giant lattice of oppositely charged ions, and explain its high melting point by the many strong electrostatic attractions that must be overcome.
• 3.6Explain in terms of structure and bonding why ionic compounds conduct electricity only when molten or dissolved, because then the ions are free to move and carry charge.
• 3.7Describe covalent bonding as the sharing of pairs of electrons between non-metal atoms, and draw dot-and-cross diagrams for simple molecules including H₂, Cl₂, O₂, N₂, HCl, H₂O, NH₃ and CH₄.
• 3.8Explain why simple molecular substances have low melting and boiling points: the intermolecular forces between molecules are weak and need little energy to overcome, while the covalent bonds inside the molecules stay intact.
• 3.9Describe how the intermolecular forces between simple molecules strengthen as molecule size increases, so bigger molecules melt and boil at higher temperatures.
• 3.10Recognise polymers as very large molecules in which thousands of atoms are joined by strong covalent bonds in long chains, with intermolecular forces between the chains, and explain why polymers are solids at room temperature.
• 3.11Recognise giant covalent structures, in which every atom is linked to its neighbours by strong covalent bonds, and explain their very high melting points and (except for graphite) their hardness and non-conductivity.
• 3.12Describe metallic bonding as positive metal ions surrounded by a sea of delocalised electrons held by strong electrostatic attraction, and use it to explain conductivity, malleability and mostly high melting points of metals.
• 3.13Explain why an alloy, containing atoms of different sizes, is harder than a pure metal: the distortion of the regular layers stops them sliding easily.
• 3.14Recognise the limitations of bonding models: dot-and-cross, 2D and 3D diagrams and ball-and-stick models each omit or exaggerate something, such as the 3D arrangement, the relative sizes of atoms and ions, or the fact that electrons are not owned by one atom.

FoundationEvery student

• 4.1Explain that carbon’s ability to form four strong covalent bonds with other carbon atoms, in chains, sheets and cages, allows it to build many different structures (allotropes) with very different properties.
• 4.2Describe diamond as a giant covalent structure in which each carbon atom bonds to four neighbours, and use this to explain its hardness, very high melting point and lack of electrical conductivity.
• 4.3Describe graphite as giant covalent layers of hexagonal rings in which each carbon bonds to three neighbours, leaving one delocalised electron per atom, and use this to explain its conductivity, slippery layers and high melting point.
• 4.4Describe graphene as a single layer of graphite, and recall its strength and electrical conductivity alongside applications such as electronics and composites.
• 4.5Recognise fullerenes as hollow cage or tube molecules of carbon (spherical fullerenes and nanotubes) and recall uses including drug delivery, lubricants, catalysts and reinforcing materials.
• 4.6Compare the properties of the carbon allotropes and relate each property to structure and bonding, including why diamond and graphite differ so sharply despite both being pure carbon.
• 4.7Recall that nanoparticles are between 1 nm and 100 nm across, and place them in the size ladder from atoms to coarse and fine particles.Chemistry only
• 4.8Calculate surface area to volume ratios and explain why shrinking a particle enormously increases the proportion of its atoms at the surface, changing how the material behaves.Chemistry only
• 4.9Describe uses of nanoparticles — in medicine (targeted drug delivery), electronics, cosmetics such as sunscreens, deodorants and as catalysts — and explain the property each use exploits.Chemistry only
• 4.10Describe why the behavior of nanoparticles may carry risks: little is known about their long-term effects, and their tiny size may let them pass into the body or the environment in ways bulk materials cannot.Chemistry only

FoundationEvery student

• 5.1Write word equations and balanced symbol equations for reactions, using chemical formulae correctly, including the diatomic elements H₂, N₂, O₂, F₂, Cl₂, Br₂ and I₂.
• 5.2Balance symbol equations by placing whole-number coefficients, understanding that balancing conserves the number of atoms of each element and never changes a formula.
• 5.3Use state symbols (s), (l), (g) and (aq) in equations and state what each represents.
• 5.4Apply the law of conservation of mass: the total mass of the products of a reaction equals the total mass of the reactants, because atoms are only rearranged.
• 5.5Explain apparent mass changes in non-enclosed systems: mass seems to fall when a gaseous product escapes and rises when a gas from the air is absorbed by the reaction mixture.
• 5.6Calculate the relative formula mass (Mr) of a compound from the relative atomic masses (Ar) of its elements.
• 5.7Deduce the empirical formula of a compound from the relative numbers of atoms present or from a model or diagram, and vice versa.
• 5.9Describe the uncertainty of a measurement: a reading is uncertain by half the resolution of the instrument, and repeats give a mean and range that describe the uncertainty of a result.

HigherHigher tier only

• 5.11Deduce an empirical formula from the masses or percentage composition of each element in the compound, dividing each mass by its relative atomic mass (Higher-tier calculation).
• 5.8Write ionic equations for reactions between ions in solution, including the cancellation of spectator ions, and balance them.
• 5.10Interpret balanced equations quantitatively: the coefficients give the ratio of moles reacting and forming, so the proportions of reactants and products can be calculated.

HigherHigher tier only

• 6.1Recall that one mole of any substance contains 6.02 × 10²³ particles (the Avogadro constant), and that the mass of one mole of a substance in grams equals its relative formula mass.
• 6.2Use the relationship moles = mass ÷ Mr to convert between the mass of a substance and the amount in moles.
• 6.3Interpret balanced symbol equations as mole ratios: the coefficients state how many moles of each substance react and form.
• 6.4Calculate the mass of a product from a given mass of reactant (and vice versa) using moles and the equation’s mole ratio.
• 6.5Identify the limiting reactant from the moles of each reactant and the equation ratio, and explain that the amount of product is set by the limiting reactant while the excess is left over.
• 6.6Use the molar gas volume: one mole of any gas occupies 24 dm³ (24 000 cm³) at room temperature and pressure, so volume of gas = moles × 24 dm³.Chemistry only
• 6.7Explain why the mole concept underpins conservation of mass: equal numbers of particles are rearranged according to the equation, and no atoms appear or disappear.

FoundationEvery student

• 7.1Use the pH scale from 0 to 14, state that acids have pH below 7, alkalis above 7 and neutral solutions exactly 7, and describe how universal indicator or a pH probe measures pH.
• 7.2Describe acids as sources of hydrogen ions, H⁺(aq), and aqueous alkalis as sources of hydroxide ions, OH⁻(aq).
• 7.3Describe neutralisation as the reaction H⁺(aq) + OH⁻(aq) → H₂O(l), forming water, with a salt as the other product.
• 7.6Describe and write equations for the reactions of acids with metals (salt + hydrogen), metal oxides and hydroxides (salt + water), and metal carbonates (salt + water + carbon dioxide).
• 7.7Predict the salt formed from an acid and a base from the acid’s family (chloride, sulfate, nitrate) and the metal ion of the base.
• 7.8Describe the required practical: preparing a pure, dry sample of a soluble salt from an insoluble base (copper oxide + sulfuric acid), including warming the acid, adding base in excess, filtering, evaporating and crystallising.
• 7.9Identify the variables, hazards and evaluation points of the salt-making practical, and explain why the base is added in excess.
• 7.11Describe uses of neutralisation including treating acidic soil with calcium hydroxide (liming) and neutralising acidic industrial waste or acid indigestion.

HigherHigher tier only

• 7.4Distinguish strong acids (fully ionised in water — HCl, HNO₃, H₂SO₄) from weak acids (partially ionised — ethanoic acid, citric acid), and concentrated from dilute solutions.
• 7.5Explain (Higher) that as pH decreases by one unit the hydrogen ion concentration increases by a factor of ten, so a pH 1 acid has 100 times the H⁺ concentration of a pH 3 acid of equal volume.
• 7.10Explain in terms of H⁺ concentration why a weak acid of the same concentration reacts more slowly and less vigorously than a strong acid.

HigherHigher tier only

• 8.1Calculate the concentration of a solution in g/dm³ from the mass of solute and the volume of solution, converting cm³ to dm³ by dividing by 1000.
• 8.2Calculate the mass of solute needed to make a solution of a given concentration and volume.
• 8.3Describe the titration technique: rinse and fill a burette with acid, measure 25.0 cm³ of alkali with a pipette into a conical flask, add a few drops of indicator, and run the acid in until the colour just changes at the endpoint.Chemistry only
• 8.4Explain the practical refinements: a rough run first, then repeats with the acid added dropwise near the endpoint, keeping only concordant titres (within 0.1 cm³) to calculate the mean.Chemistry only
• 8.5Calculate the concentration of a solution in mol/dm³ (moles of solute ÷ volume in dm³) and convert between mol/dm³ and g/dm³ via the Mr.Chemistry only
• 8.6Carry out titration calculations: use moles = concentration × volume to find the moles reacting, apply the equation ratio, and work back to the unknown concentration.Chemistry only
• 8.7Explain why titrations are used to make salts from soluble reactants (such as ammonium chloride) where no excess insoluble solid can be filtered off, and why no indicator is used when preparing a salt this way.Chemistry only
• 8.8Describe the uncertainty of a burette reading (±0.05 cm³) and how repeating titrations and averaging concordant values reduces the uncertainty of the result.Chemistry only

FoundationEvery student

• 9.1State the order of the reactivity series (potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold) and explain that it ranks metals by how easily their atoms form positive ions.
• 9.2Describe how reactions with water and dilute acids place metals in the series, and recall that metals below hydrogen do not react with dilute acids.
• 9.3Use displacement reactions between metals and metal salt solutions to compare reactivity, predicting whether a reaction occurs.
• 9.4Define oxidation as the gain of oxygen and reduction as the loss of oxygen, identifying each in reactions such as metal extraction with carbon.
• 9.6Explain why metals less reactive than carbon (zinc, iron, copper) are extracted from their oxides by reduction with carbon, and why metals more reactive than carbon (aluminium and above) need electrolysis.
• 9.7Describe the extraction of iron in the blast furnace in outline: hot air, coke burning to carbon dioxide and carbon monoxide, reduction of iron(III) oxide, and molten iron tapped at the base.
• 9.9Explain why a metal’s method of extraction drives its cost, and why gold’s unreactivity lets it occur native while reactive metals are only ever found as compounds.

HigherHigher tier only

• 9.5Explain (Higher) oxidation and reduction in terms of electron loss and gain (OIL RIG), and write ionic equations and half equations for displacement reactions.
• 9.8Describe (Chemistry only) the biological extraction methods: phytoextraction (plants absorbing metal ions, then being burned to leave ash rich in metal compounds) and bioleaching (bacteria producing leachate containing metal ions), and evaluate their advantages of low energy and use of low-grade ores against their slowness.Chemistry only

FoundationEvery student

• 10.1Define electrolysis as splitting a molten or dissolved ionic compound using electricity, and explain why the compound must be molten or dissolved — the ions must be free to move.
• 10.2Identify the electrodes: the negative cathode attracts positive ions (cations), and the positive anode attracts negative ions (anions).
• 10.3Predict the products of electrolysis of a molten binary ionic compound: the metal forms at the cathode and the non-metal at the anode.
• 10.4Predict the products of aqueous electrolysis: at the cathode hydrogen forms unless the metal is less reactive than hydrogen (copper, silver); at the anode a halide forms halogen if present, otherwise oxygen from hydroxide ions.
• 10.5Describe the extraction of aluminium: molten aluminium oxide with cryolite added to lower the melting point, aluminium at the cathode, oxygen at the anode, and the carbon anodes burning away.
• 10.8Describe the required practical: electrolysis of aqueous solutions (sodium chloride, copper(II) sulfate with inert electrodes, dilute sulfuric acid) using inert electrodes, identifying the products at each electrode with the gas tests.

HigherHigher tier only

• 10.6Write half equations for the reactions at each electrode (Higher), balancing atoms and charge, such as Al³⁺ + 3e⁻ → Al and 2Cl⁻ → Cl₂ + 2e⁻.
• 10.7Explain oxidation and reduction at each electrode in terms of electrons: reduction (gain) at the cathode, oxidation (loss) at the anode.
• 10.9Explain observations of the aqueous electrolysis practical: the 2 : 1 hydrogen-to-oxygen volume ratio from dilute sulfuric acid, chlorine’s bleaching effect from concentrated brine, and copper depositing on the cathode.

FoundationEvery student

• 11.1Classify reactions as exothermic (temperature of surroundings rises) or endothermic (temperature falls), giving examples including combustion, neutralisation and oxidation as exothermic, and thermal decomposition and citric acid with sodium hydrogencarbonate as endothermic.
• 11.2Describe the required practical: measuring temperature changes when solutions react, including the method, variables, plotting a graph of temperature change against volume and using it to find the point of exact neutralisation.
• 11.3Draw and interpret reaction profiles for exothermic and endothermic reactions, labelling reactants, products, the overall energy change (ΔH) and the activation energy.
• 11.4Explain that activation energy is the minimum energy particles need to react, and that catalysts lower it.
• 11.6Describe (Chemistry only) how a chemical cell makes electricity from a redox reaction, how voltage depends on the metals and electrolyte, and that cells go flat when their reactants run out.Chemistry only
• 11.7Describe (Chemistry only) the hydrogen–oxygen fuel cell: hydrogen at the anode and oxygen at the cathode producing clean water, with overall equation 2H₂ + O₂ → 2H₂O, and write its half equations.Chemistry only
• 11.8Evaluate (Chemistry only) fuel cells against rechargeable cells: fuel cells need a continuous fuel supply and cost more, but never need recharging and produce no greenhouse gases at the point of use.Chemistry only

HigherHigher tier only

• 11.5Explain exothermic and endothermic changes (Higher) in terms of bond breaking (energy taken in) being greater or smaller than bond making (energy released), and calculate ΔH using bond energies.

FoundationEvery student

• 12.1Describe practical methods for measuring the rate of a reaction: collecting gas in a syringe, measuring mass loss on a balance, and timing the appearance of a precipitate that obscures a mark.
• 12.2Calculate the mean rate from measured quantities, and find the rate at a specific time from the gradient of a tangent on a rate graph.
• 12.3Explain rates using collision theory: reactions occur when particles collide with at least the activation energy, and rate depends on collision frequency and energy.
• 12.4Explain the effects of temperature, concentration/pressure, and surface area on rate using collision theory.
• 12.5Describe catalysts as substances that speed reactions without being used up, and explain that they lower the activation energy by providing an alternative pathway.
• 12.6Describe the required practical: investigating how concentration affects the rate (magnesium with hydrochloric acid collecting hydrogen, or sodium thiosulfate with hydrochloric acid obscuring a cross), including variables, method and evaluation.
• 12.7Interpret rate graphs: steeper gradient means faster rate, the graph flattens as reactants are used up, and tangents give instantaneous rates.
• 12.8Explain why catalysts are important industrially (less energy, lower cost, fewer emissions) and why enzymes are biological catalysts.

FoundationEvery student

• 13.1Describe reversible reactions and dynamic equilibrium: in a closed system, when the forward and reverse reactions happen at the same rate, the concentrations of reactants and products stay constant.
• 13.4Describe the Haber process: purified nitrogen and hydrogen passed over an iron catalyst at about 450 °C and about 200 atmospheres, ammonia condensed out as a liquid and unreacted gases recycled.
• 13.5Calculate the percentage yield of a reaction from actual and theoretical yields, and explain why a reversible reaction never converts all reactants to product.Chemistry only
• 13.6Define and calculate (Chemistry only) the atom economy of a reaction: the percentage of the mass of the products that is the desired product; compare routes and explain why industry prefers high atom economy, including finding uses for co-products.Chemistry only
• 13.7Describe (Chemistry only) NPK fertilisers as formulations of nitrogen, phosphorus and potassium compounds; explain how ammonium salts are made by neutralising acids with ammonia and how phosphate rock is treated with nitric, sulfuric or phosphoric acid to give soluble salts.Chemistry only
• 13.8Compare (Chemistry only) making a fertiliser salt in a school laboratory (small-scale batch, slower, dilute solutions) with industrial production (large-scale continuous, faster per tonne), and evaluate the choices.Chemistry only

HigherHigher tier only

• 13.2State (Higher) Le Chatelier’s principle: when a change is made to the conditions of a system at equilibrium, the position of equilibrium shifts to oppose the change.
• 13.3Predict (Higher) the effect of changing concentration, pressure (for gases) or temperature on the position of equilibrium, and explain why a catalyst changes the rate but not the position.

FoundationEvery student

• 14.1Describe crude oil as a finite resource found in rocks, a complex mixture of hydrocarbons formed over millions of years from the remains of ancient biomass (mainly plankton) buried in mud.
• 14.2Define hydrocarbons as compounds of hydrogen and carbon only, and describe alkanes as saturated hydrocarbons with the general formula CnH₂n+₂, knowing the first four (methane, ethane, propane, butane).
• 14.3Describe fractional distillation: crude oil is heated and vaporised, then fractions condense at different heights in the column according to their boiling points.
• 14.4Describe the trends in the properties of fractions with chain length — viscosity, flammability and boiling point — and how this affects their uses as fuels and feedstocks.
• 14.5Write balanced equations for the complete combustion of hydrocarbons, and describe the products of incomplete combustion (carbon monoxide and carbon/soot) with their hazards.
• 14.6Describe cracking as a thermal decomposition reaction splitting long alkanes into shorter, more useful alkanes and alkenes, using high temperatures and a catalyst or steam.
• 14.7Describe alkenes as unsaturated hydrocarbons with the general formula CnH₂n containing a carbon–carbon double bond, and describe the bromine water test that distinguishes them from alkanes.
• 14.8Explain why cracking is needed commercially: the supply of long-chain fractions from distillation exceeds demand, while short-chain hydrocarbons (fuels and polymer feedstocks) are in high demand.

FoundationEvery student

• 15.1Describe the alcohols as a homologous series with the functional group –OH, knowing methanol, ethanol and propanol, and that ethanol can be made by fermentation of glucose or by reacting ethene with steam.Chemistry only
• 15.2Describe the combustion of alcohols in oxygen releasing energy, and the oxidation of ethanol to ethanoic acid (including microbial oxidation causing wine to turn to vinegar).Chemistry only
• 15.3Describe aqueous carboxylic acids as weak acids: they are only partially ionised in solution, so they react more slowly than strong acids of the same concentration — with metals, bases and carbonates.Chemistry only
• 15.4Describe addition polymerisation: small alkene monomers joining with no other product, and draw or interpret the repeat unit of polymers such as poly(ethene).
• 15.6Describe (Chemistry only) condensation polymerisation: monomers with two functional groups joining with loss of a small molecule such as water, forming polyesters or polypeptides.Chemistry only
• 15.7Describe (Chemistry only) the structures of DNA and natural polymers: amino acids as monomers with –NH₂ and –COOH groups forming polypeptides and proteins, and DNA as two polymer strands of nucleotides cross-linked in a double helix.Chemistry only

HigherHigher tier only

• 15.5Explain (Higher) how the properties of addition polymers depend on the monomer and how intermolecular forces between chains affect strength and melting behaviour.
• 15.8Compare (Higher, Chemistry only) addition and condensation polymerisation: number of monomer types, the other product formed, and the bonds broken or made in each.Chemistry only

FoundationEvery student

• 16.1Describe a pure substance as a single element or compound (not a mixture), and explain that melting or boiling points can be used to distinguish pure from impure substances.
• 16.2Describe formulations as mixtures designed for a specific purpose, each component present in a measured amount, with examples such as fuels, alloys, fertilisers and medicines.
• 16.3Describe paper chromatography: the mobile phase (solvent) carries dissolved substances up the stationary phase (paper), separating them by their different attractions to each phase.
• 16.4Calculate Rf values (distance moved by substance ÷ distance moved by solvent), use them to identify substances against references, and interpret chromatograms to judge purity.
• 16.5Describe the tests for gases: hydrogen (burning splint gives a squeaky pop), oxygen (relights a glowing splint), carbon dioxide (turns limewater cloudy) and chlorine (bleaches damp litmus paper white).
• 16.6Describe (Chemistry only) the flame test colours: lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper green.Chemistry only
• 16.7Describe (Chemistry only) the sodium hydroxide precipitate tests for metal ions and the tests for anions: carbonates (dilute acid gives CO₂), sulfates (acidify then barium chloride gives white precipitate) and halides (acidify then silver nitrate gives white, cream or yellow precipitates).Chemistry only
• 16.8Explain (Chemistry only) why instrumental methods such as flame emission spectroscopy are used to identify elements and metal ions: they are fast, sensitive, accurate and can analyse mixtures, using calibration data.Chemistry only

FoundationEvery student

• 17.1Describe the theories for the evolution of the atmosphere in phases: volcanic activity releasing carbon dioxide, water vapour and nitrogen; the oceans forming as water vapour condensed and dissolving much of the carbon dioxide; and green plants and algae increasing oxygen through photosynthesis.
• 17.2Describe today’s atmosphere as approximately 78% nitrogen, 21% oxygen, 1% argon and about 0.04% carbon dioxide, and how these proportions changed since the early atmosphere.
• 17.3Describe the greenhouse effect: greenhouse gases (water vapour, carbon dioxide, methane) absorb outgoing long-wave radiation and re-radiate it, keeping Earth warmer than it would otherwise be.
• 17.4Explain the enhanced greenhouse effect: human activities increasing greenhouse gas levels cause more energy to be trapped, raising Earth’s temperature (global warming).
• 17.5Describe the evidence for human-caused climate change: peer-reviewed evidence based on systematic measurements and multiple independent sources (historical records, ice cores, tree rings, temperature records).
• 17.6Define carbon footprint as the total greenhouse gases emitted over the full life cycle of a product, service or event; describe ways to reduce it and why reductions are difficult (international cooperation, economic costs).
• 17.7Describe the atmospheric pollutants from combustion: carbon monoxide (toxic, formed by incomplete combustion), soot/particulates (health problems, global dimming), sulfur dioxide and oxides of nitrogen (acid rain, breathing problems).
• 17.8Explain how the pollutants form: sulfur dioxide from burning sulfur impurities in fuels, and oxides of nitrogen from nitrogen and oxygen reacting at the high temperatures inside engines.

FoundationEvery student

• 18.1State that finite resources (fossil fuels, metal ores, minerals) are used faster than they form, describe the difference with renewable resources, and explain what sustainable development means.
• 18.2Describe potable water as water that is safe to drink (not chemically pure), and describe how UK water is treated: choosing a fresh source, sedimentation and filtration to remove solids, and sterilisation with chlorine, ozone or ultraviolet light to kill microbes.
• 18.3Describe desalination of seawater by distillation or reverse osmosis, and evaluate the energy costs of producing potable water in different contexts.
• 18.4Describe the required practical: analysing a water sample — measuring its pH, evaporating a measured volume to find the mass of dissolved solids, and distilling the sample to obtain pure water.
• 18.5Describe how waste water is treated: screening, sedimentation into sludge and effluent, anaerobic digestion of sludge (producing biogas), and aerobic biological treatment of effluent before safe release.
• 18.6Describe a life cycle assessment: the stages from extracting raw materials through manufacturing, use and disposal, listing impacts at each stage, and explain why LCAs can be subjective and selective.
• 18.7Describe ways of reducing the use of resources: recycling (metals, glass, plastics) and reuse, including the energy savings versus extracting and processing raw materials.
• 18.8Describe (Chemistry only) corrosion and its prevention: rusting as iron reacting with oxygen and water to form hydrated iron(III) oxide, prevented by barrier methods and sacrificial protection.Chemistry only
• 18.9Describe (Chemistry only) alloys: gold alloys and carats, bronze (copper–tin), brass (copper–zinc), and steels of different carbon content, explaining how alloying changes properties.Chemistry only
• 18.10Describe (Chemistry only) the properties and uses of ceramics, polymers and composites: pottery and glass, thermosoftening versus thermosetting polymers, and composites as a matrix reinforced by fibres or particles.Chemistry only